Reduction potential

Redox potential (also known as oxidation / reduction potential, ORP, pe, <math>E_{red}</math>, or <math>E_{h}</math>) is a measure of the tendency of a chemical species to acquire electrons from, or lose electrons to, an electrode and thereby be reduced or oxidised respectively. Redox potential is expressed in volts (V). Each species has its own intrinsic redox potential; for example, the more positive the reduction potential (reduction potential is more often used due to general formalism in electrochemistry), the greater the species' affinity for electrons and tendency to be reduced.

Measurement and interpretation

In aqueous solutions, redox potential is a measure of the tendency of the solution to either gain or lose electrons in a reaction. A solution where the solute has a higher (more positive) reduction potential than some other molecule will have a tendency to gain electrons from this molecule (i.e. to be reduced by oxidizing this other molecule), and a solution with a lower (more negative) reduction potential will have a tendency to lose electrons to other substances (i.e. to be oxidized by reducing the other substance). Because the absolute potentials are almost impossible to accurately measure, reduction potentials are defined relative to a reference electrode. Reduction potentials of aqueous solutions are determined by measuring the potential difference between an inert sensing electrode in contact with the solution and a stable reference electrode connected to the solution by a salt bridge.

The sensing electrode acts as a platform for electron transfer to or from the reference half cell; it is typically made of platinum, although gold and graphite can be used as well. The reference half cell consists of a redox standard of known potential. The standard hydrogen electrode (SHE) is the reference from which all standard redox potentials are determined, and has been assigned an arbitrary half cell potential of 0.0 V. However, it is fragile and impractical for routine laboratory use. Therefore, more stable reference electrodes, such as silver chloride and saturated calomel (SCE), are commonly used because of their more reliable performance.

Although measurement of the redox potential in aqueous solutions is relatively straightforward, many factors limit its interpretation, such as effects of solution temperature and pH, irreversible reactions, slow electrode kinetics, non-equilibrium, presence of multiple redox couples, electrode poisoning, small exchange currents, and inert redox couples. Consequently, practical measurements seldom correlate with calculated values. Nevertheless, reduction potential measurement has proven useful as an analytical tool in monitoring changes in a system rather than determining their absolute value (e.g. process control and titrations).

Explanation

Similar to how the concentration of hydrogen ions determines the acidity or pH of an aqueous solution, the tendency of electron transfer between a chemical species and an electrode determines the redox potential of an electrode couple. Like pH, redox potential represents how easily electrons are transferred to or from species in solution. Redox potential characterises the ability under the specific condition of a chemical species to lose or gain electrons instead of the amount of electrons available for oxidation or reduction.

The notion of is used with Pourbaix diagrams. is a dimensionless number and can easily be related to EH by the following relationship:

: <math>pe = \frac{E_{H{V_T \lambda} = \frac{E_{H{0.05916} = 16.903 \, \text{×} \, E_{H}</math>

where, <math>V_T=\frac{RT}{F}</math> is the thermal voltage, with , the gas constant (), , the absolute temperature in Kelvin (298.15 K = 25 °C = 77 °F), , the Faraday constant (96 485 coulomb/mol of ), and λ = ln(10) ≈ 2.3026.

In fact, <math>pe = -\log[e^-]</math> is defined as the negative logarithm of the free electron concentration in solution, and is directly proportional to the redox potential. Sometimes <math>pe</math> is used as a unit of reduction potential instead of <math>E_h</math>, for example, in environmental chemistry.

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: 3 + = + 2 + 2

Attention, I have not changed the numerical values, nor any sign in the corresponding Nernst equation as I have no access to the book of Garrels and Christ (1990) -->

where:

The slope 0.0296 of the line is −1/2 of the −0.05916 value above, since . Additionally, the value −0.0885 corresponds to −0.05916 × 3/2.

Biochemistry

Many enzymatic reactions are oxidation–reduction reactions, in which one compound is oxidized and another compound is reduced. The ability of an organism to carry out oxidation–reduction reactions depends on the oxidation–reduction state of the environment, or its reduction potential (<math>E_h</math>).

Strictly aerobic microorganisms are generally active at positive <math>E_h</math> values, whereas strict anaerobes are generally active at negative <math>E_h</math> values. Redox affects the solubility of nutrients, especially metal ions.

There are organisms that can adjust their metabolism to their environment, such as facultative anaerobes. Facultative anaerobes can be active at positive Eh values, and at negative Eh values in the presence of oxygen-bearing inorganic compounds, such as nitrates and sulfates.

In biochemistry, apparent standard reduction potentials, or formal potentials, (<math>E^{\ominus '}_{\text{red</math>, noted with a prime mark in superscript) calculated at pH 7, closer to the pH of biological and intra-cellular fluids, are used to more easily assess if a given biochemical redox reaction is possible. They must not be confused with the common standard reduction potentials </math>) determined under standard conditions (; ) with the concentration of each dissolved species being taken as 1 M, and thus .

Environmental chemistry

In the field of environmental chemistry, the reduction potential is used to determine if oxidizing or reducing conditions are prevalent in water or soil, and to predict the states of different chemical species in the water, such as dissolved metals. pe values in water range from −12 to 25; the levels where the water itself becomes reduced or oxidized, respectively.

Water quality

The oxido-reduction potential (ORP) can be used for the systems monitoring water quality with the advantage of a single-value measure for the disinfection potential, showing the effective activity of the disinfectant rather than the applied dose. For example, E. coli, Salmonella, Listeria, and other pathogens have survival times of less than 30 seconds when the ORP is above 665 mV, compared to more than 300 seconds when ORP is below 485 mV.

Geochemistry and mineralogy

Eh–pH (Pourbaix) diagrams are commonly used in mining and geology for assessment of the stability fields of minerals and dissolved species. Under the conditions where a mineral (solid) phase is predicted to be the most stable form of an element, these diagrams show that mineral. As the predicted results are all from thermodynamic (at equilibrium state) evaluations, these diagrams should be used with caution. Although the formation of a mineral or its dissolution may be predicted to occur under a set of conditions, the process may practically be negligible because its rate is too slow. Consequently, kinetic evaluations at the same time are necessary. Nevertheless, the equilibrium conditions can be used to evaluate the direction of spontaneous changes and the magnitude of the driving force behind them.

Reduction potential

Measurement and interpretation

Explanation

Biochemistry

Environmental chemistry

Water quality

Geochemistry and mineralogy

See also

References

Notes

Additional notes

External links